Transcript Notes -- Unit 5 -- Reactions and Stoichiometry

Stoichiometry & Solution Stoich
Chapter 3 and 4.2/4.6
Chemical Reactions/Equations
 Reactants
are listed on the left
 Products on the right
 Atoms are neither created nor
destroyed
 Shows what state of matter the
compounds are in: Solid (s), Liquid (l),
Gas (g), or dissolved in water (aqueous
solution) (aq)
Chemical Equations

Hydrogen burns and reacts with oxygen in
the air to form water. write the chemical
reaction:
H2 (g)+ O2 (g)  H2O (l)
 What’s wrong with the above equation?
 Due to the Law of Conservation of
Mass, it balances out to:
2H2 (g)+ O2 (g) 2H2O (l)
Balancing Equations
 Determine
what reaction is occurring
(sometimes it helps to write it in word
form)
 Write the unbalanced (skeleton) eqn.
 Balance the equation by inspection, by
adding coefficients (usually works best
going from left to right)
 Include phase information
Balancing Equations

Al2(SO4)3 + Ca(OH)2  Al(OH)3 + CaSO4
 H3PO4
+ NaOH  Na3PO4 + HOH
Types of Reactions
 Knowing
the basic types of chemical
reactions, helps you to predict the
products of many reactions
 Combination Reaction:
2Mg (s) + O2 (g)  2MgO (s)
 Decomposition Reaction:
CaCO3 (s)  CaO (s) + CO2 (g)
 Combustion
Reactions:
 Rapid reactions that produce a flame
 Most involve O2 (from air) as a reactant
 When Hydrocarbons (CxHy) react with
O2 they produce CO2 and H2O.
C3H8(g) + 5O2 (g)  3CO2 (g) + 4H2O(g)

When an air bag deploys,
sodium azide (NaN3)
decomposes, rapidly
releasing nitrogen gas and
sodium.
1. What type of RXN took
place?
2. Write a balanced
equation
Balancing Combustion RXNs
 first
start with those elements that occur
in the fewest chemical formulas.
 Try a few:
CH4 + O2  CO2 + H2O
C2H5OH + O2  CO2 + H2O
How much sand?
3 Ways to Measure
Matter
COUNT – 1 million grains of sand
 By MASS – 1,000 grams of sand
 By VOLUME – 100 liters of sand
 By
ATOMIC MASS:
What is the atomic mass of Hydrogen?
 1.01 a.m.u.
What is the atomic mass of Oxygen?
 15.999  16.0 a.m.u.
=
SO3
+
1 S atom
3 O atoms
You can calculate the mass of a molecule by
adding the atomic masses of the atoms making up
the molecule. (32.1 + 16 + 16 + 16 = 80.1 amu)
Formula Weights
What is the atomic mass of Water (H2O)?
 2H – 2 (1.0) = 2.0 a.m.u.
 1O – 1(16.0) =16.0 a.m.u.
18.0 a.m.u
What is the atomic mass of Ca(NO3)2?
 1 Ca – 1(40.1) = 40.1 a.m.u
 2 N – 2 (14.0) = 28.0 a.m.u.
 6 O – 6 (16.0) = 96 a.m.u.
164.1 a.m.u.
Percentage Composition
 The
percentage by mass contributed by
each element in the substance
 Calculating any percentage is just like
calculating your grade in a class (“the
part, divided by the whole, multiplied
by 100”)
 Calculate the percentage by mass of
each element in Ca(NO3)2
 Calculate
the percentage composition
of Oxygen in glucose (C6H12O6)
Problem Solving Hints
1.
2.
3.
4.
Analyze the Problem
Develop a plan for solving the problem
Solve the problem
Check the solution
The Mole
What is a MOLE?
is a quantity equal to Avogadro’s
Number (6.02 x 1023)
 6.02 x 1023 particles (atoms or molecules)
– depending on what you are looking at.
 A mole of anything contains the same
number of “things” as a mole of anything
else.
 one mole is set by defining one mole of carbon
12 atoms to have a mass of exactly 12 grams.
 A mole
Molar Mass
 The
mass in grams of one mole of a
compound
 It is numerical equivalent to what it’s
atomic weight was in a.m.u.’s
Molar Mass of Water (H2O)?
 18 grams/mole
Molar Mass of Calcium Nitrate?
 161.4 grams/mole
divide
Multiply
by 22.4
multiply
Divide by
22.4
Use 6.02 x 1023
Practice Problems
 How
many moles are in 5g of Copper?
 How many grams in 3.2 moles of Oxygen?
 How many atoms in .350 moles of
Sodium?
 How many moles in 7.2 x 1024 atoms of
gold?
 How many glucose molecules in 5.23
grams of glucose, C6H12O6?
 How many atoms of Oxygen?
Determining the Empirical
Formula of a Compound
1.
2.
3.
4.
Determine the percentage of each element
in your compound
Treat % as grams, and convert grams of
each to moles of each element
Find Smallest whole number ratio (divide
the larger number by the smaller one)
If ratio is not all whole numbers, multiply
each by an integer to make all whole
number ratio
Example 1

1.
2.
3.
Mercury forms a compound with
chlorine that is 73.9% mercury and
26.1% chlorine.
Convert to 73.9 g Hg and 26.1 g Cl
Convert to moles of each element
Find Smallest whole number ratio
(divide the larger number by the
smaller one)
Example 2
 Ascorbic
acid (Vitamin C) contains
40.92% C, 4.58% H, and 54.5% O by
mass. What is the empirical formula of
ascorbic acid?
C:H:O = 3 (1:1.33:1) = 3:4:3
C3H4O3
Determining the molecular
Formula
 Find
the empirical formula mass.
 Divide the known molecular mass by
the empirical formula mass, deriving a
whole number, n.
 Multiply the empirical formula by n to
derive the molecular formula.
Example 1

1.
2.
3.
The empirical formula of ascorbic acid is
C3H4O3
The empirical formula mass is 88.0 amu.
The experimentally determined
molecular weight is 176 amu.
Therefore the molecule has twice the
mass (176/88 = 2.00) and must have
twice as many of each atom for a
molecular formula of C6H8O6
Example 2

Mesitylene, has an empirical formula of
C3H4. The experimentally determined
molecular weight of this substance is 121
amu. What is the molecular formula of
mesitylene?
Combustion Analysis to
determine empirical formulas
 Used
to figure out how much Carbon
and Hydrogen were in the original
sample of the Hydrocarbon
Example 1

1.
Isopropyl alcohol (rubbing alcohol) is
composed of C, H, and O. Combustion of
.255g of isopropyl alcohol produces 0.561
g CO2 and 0.306g H2O. Determine the
empirical formula of isopropyl alcohol.
Calculate the number of grams of C
present in the CO2 and grams of H present
in H2O using the mole concept and
dimensional analysis.
2.
3.
4.
Calculate the mass of O in the final
sample by subtracting the mass of C
and H in the sample from the total
sample mass: Mass of O = mass of
sample – (mass of C + mass of H)
Then calculate the number of moles of
C, H, and O in the sample
Find the lowest whole number ratio
Example 2
 Caproic
acid, which is responsible for the
foul odor of dirty socks, is composed of C,
H, and O atoms. Combustion of a 0.225g
sample of this compound produces 0.512g
CO2 and 0.209 g H2O. What is the
empirical formula of caproic acid?
 It has a molar mass of 116 g/mol. What is
its molecular formula?
Stoichiometric Calculations:
Quantitative info from balanced EQNs
1.
2.
3.
4.
Balance the chemical equation
Convert grams of reactant or product
to moles.
Compare moles of the known to moles
of the desired substance (use a ratio
derived from the coefficients in the
balanced equation.)
Convert from moles back to grams if
required.
Chemical Calculations
 Mole
 Mole
to Mole
to Gram
m
m
m
m
g
 Gram
to Gram
m
m
g
g
m
m
g
g
Mole-Mole Calculations
 How
many moles of ammonia (NH3)
are produced when 0.60 moles of
nitrogen reacts with hydrogen?
 Step 1: Write Chemical EQN
 N 2 + H2 
NH3
 Step 2: Balance EQN
 N2 + 3H2  2NH3
 Step
3: Find known & unknown,
then calculate
Known: 0.60 mol N2
Unknown: ? mol NH3
Mole Ratio: 1 mol N2
2 mol NH3
0.60 mol N2 x 2 mol NH3 = 1.2 mol NH3
1 mol N2
Gram-Gram Calculations
 Calculate
the number of grams of NH3
produced by the reaction of 5.40 g of
hydrogen with an excess of nitrogen.
m
m
m
g
m
g
N2 + 3H2
5.40 g
1 g
2
2NH3
? grams
g 3
Step 1 : Change grams to moles.
5.40 g H2 x 1 mol H2 = 2.70 mol H2
2.0 g H2
Step 2 : Get a mole ratio from
the equation.
3 mol H2
2 mol NH3
then
2.7 mol H2
? mol NH3
Step 3 : Solve for the unknown
number of moles.
(2.7 x 2) / 3 = 1.8 moles NH3
Step 4 : Change moles of the
unknown to grams.
1.8 mole NH3 x 17.0 g NH3 = 30.6 g NH3
1 mol NH3
Examples
 How
many grams of water are
produced in the oxidation of 1.00 g of
glucose, C6H12O6?
 Propane, C3H8, is a common fuel used
for cooking and home heating. What
mass of O2 is consumed in the
combustion of 2.75 moles of propane?
Examples
+ NH3(aq) 
Pt(NH3)2Cl2 (s)+ KCl(aq)
 what mass of Pt(NH3)2Cl2 can be
produced from 65 g of K2PtCl4 ?
 How much KCl will be produced?
 How much from 65 grams of NH3?
 K2PtCl4(aq)
Making Chocolate Chip Cookies
Ingredients in Kitchen (I have a BIG kitchen):
 40 lbs of butter
 2 lbs of salt
 1 gallon of vanilla extract
 80 lbs of chocolate chips
What’s going
 200 lbs of flour
to determine
 150 lbs of sugar
how many cookies
 10 lbs baking soda
I can make?
 2 eggs
Limiting Reactants
 limits
or determines the amount of
product that can be formed in a
reaction.
 The reactant that isn’t used up is called
the excess reagent
 To determine, book says use ratio
method and I.C.E. chart, I’ll show you a
different method… both work.
 To
determine the limiting reagent requires
that you do two stoichiometry problems.
 Figure out how much product each
reactant makes.
 The one that makes the least is the
limiting reagent
Using an I.C.E. chart

Convert both reactants into moles… suppose
we had 10 moles of H2 and 7 moles of O2.
2H2 + O2  2H2O
Initial
Quantities
Change
(reaction)
Expected
Quantities
10 mol
7 mol 0 mol
-10 mol -5 mol +10 mol
0 mol
2 mol
10 mol
Example
 Ammonia
is produced by the following
reaction
N2 + H2  NH3
 What mass of ammonia can be
produced from a mixture of 100. g N2
and 500. g H2 ?
 How much unreacted material
remains?
Success of Reaction
 The
amount of stuff you make is the
yield.
 The theoretical yield is the amount you
would make if everything went perfect.
 The actual yield is what you make in the
lab.
Percent Yield
%
yield = Actual
x 100%
Theoretical
%
yield = what you got
x 100%
what you could have got
Examples
 Aluminum
burns in bromine producing
aluminum bromide. In a laboratory 6.0 g
of aluminum reacts with excess
bromine. 50.3 g of aluminum bromide
are produced. What are the three types
of yield. (actual, theoretical, percent)
Precipitation Reactions
 Occur
when certain pairs of oppositely
charged ions attract each other so
strongly that they form an insoluble
ionic solid.
 Can determine if a precipitate will form
by following certain guidlines
Solubility Guidelines for
Ionic Compounds
Previous picture: Only 1.2 x10-3 mol of PbI2
dissolves in a liter of water.
 We will consider it insoluble if the solubility is
less than 0.01 mol/L
 Refer to Table 4.1 (and Ksp values on pg 1045)
in text for solubility guidelines for common ionic
compounds in water
 Dateless Dudes:
Ammonium, Nitrate, Acetate, and Alkali Metals

 Ksp
= Solubility-Product Constants
pg. 1045 in text book
 The smaller the number, the more likely it
will precipitate
 If there is no insoluble product, the
reaction does not occur.
 Exchange (Metathesis) Reactions are
another name for double replacement
reactions… (+ and - ions switch partners)
Simple Solubility Rules
1.
2.
3.
Most nitrate slats are soluble
Most salts containing the alkali metal
ions and ammonium are soluble
Most chloride, bromide, and iodide
salts are soluble. (exceptions are
salts containing the ions Ag+, Pb+2,
and Hg+2)
4.
5.
6.
Most sulfate (SO4-2) salts are soluble.
(notable exceptions are BaSO4,
PbSO4, HgSO4, and CaSO4)
Most OH- salts are only slightly
soluble. The important soluble ones
are NaOH and KOH. Ba, Sr and Ca
are marginally soluble.
Most S-2, CO3-2, CrO4-2 and PO4-3 are
slightly soluble.
Ionic and Net Ionic Equations
1.
Write out the molecular equation:
Pb(NO3)2(aq)+ 2 KI(aq)  PbI2 (s) + 2KNO3(aq)
2.
3.
Write out the complete ionic equation
and Identify and cancel out spectator
ions:
Write out the Net Ionic equation:
Example
An aqueous solution of silver nitrate
reacts with an aqueous solution of
potassium chloride. A precipitate is
produced.
1. Determine the precipitate using Ksp.
2. Write the Molecular, Ionic, and Net
Ionic equations for the above reaction
Solution Stoichiometry and
Chemical Analysis
 Molarity
= moles of solute / Liters of
solution.
 M1V1 = M2V2
 We can take the molarity and volume of
a solution to find out the moles of a
solution. And then use dimensional
analysis to determine the moles or
grams of another reactant or product.
Example 1
How many grams of Ca(OH)2 are
needed to neutralize 25.0 mL of 0.100
M HNO3?
1. Use molarity and volume of HNO3 to
convert to moles of HNO3
 Balance equation:
HNO3 + Ca(OH)2  H2O + Ca(NO3)2
1. Convert moles of HNO3 to moles of
Ca(OH)2 and then to grams of Ca(OH)2

Example 2
 How
many grams of NaOH are needed
to neutralize 20.0 mL of 0.150M H2SO4
solution?
 NaOH + H2SO4  H2O + Na2SO4